Matter and Energy Notes

Classification of Matter

• Matter:

  • Material that makes up all things.

  • Has mass and occupies space.

• Classification of Matter:

  • Pure Substances: Fixed or definite composition.

  • Mixtures: Two or more different substances physically mixed, not chemically combined.

Pure Substances: Elements and Compounds

• Pure Substance:

  • Type of matter with a fixed or definite composition.

  • Element: Composed of one type of atom (e.g., copper, Cu; lead, Pb; aluminum, Al).

    • There are 117 known elements, 88 of which are naturally occurring.

  • Compound: Composed of two or more elements always combined in the same proportion (e.g., hydrogen peroxide, H2O2; table salt, NaCl; sugar, C{12}H{22}O{11}; water, H2O).

Mixtures

• Mixture:

  • Two or more substances physically mixed, not chemically combined.

  • Substances in different proportions.

  • Separation possible via physical methods (e.g., filtration).

• Homogeneous Mixtures:

  • Uniform composition throughout.

  • Different parts not visible (e.g., brass - copper and zinc).

  • Scuba breathing mixtures (nitrox, heliox, trimix).

• Heterogeneous Mixtures:

  • Composition varies from one part to another.

  • Different parts are visible (e.g., copper metal and water).

States and Properties of Matter

States of Matter

• Solids:

  • Definite shape and volume.

  • Particles close together in a fixed arrangement.

  • Particles move very slowly.

• Liquids:

  • Indefinite shape, definite volume.

  • Take the shape of their container.

  • Particles close together but mobile.

  • Particles move slowly.

• Gases:

  • Indefinite shape and volume.

  • Take the shape and volume of their container.

  • Particles far apart.

  • Particles move very fast.

Properties of Matter

• Physical Properties:

  • Observed/measured without changing the substance's identity.

  • Include shape, physical state, boiling/freezing points, density, color.

  • Example: Copper (reddish-orange, shiny, conducts heat/electricity, solid at 25 °C, melting point 1083 °C, boiling point 2567 °C).

• Physical Change:

  • Change in state or physical shape.

  • No change in the identity/composition of the substance.

• Chemical Properties:

  • Describe a substance's ability to interact with other substances and change into new substances.

• Chemical Change:

  • Original substance transforms into a new substance with new chemical and physical properties.

  • Example: Sugar caramelizing.

Temperature

• Temperature:

  • Measure of hotness or coldness of an object compared to another.

  • Indicates heat flow (higher to lower temperature).

  • Measured using a thermometer.

Temperature Scales

• Fahrenheit, Celsius, and Kelvin.

• Reference points: boiling and freezing points of water.

• Fahrenheit vs. Celsius:

  • 100 °C between freezing and boiling points of water on Celsius scale.

  • 180 °F between freezing and boiling points of water on Fahrenheit scale.

  • 180
    ewline Fahrenheit
    ewline degrees = 100
    ewline degrees
    ewline Celsius

Converting Between Celsius and Fahrenheit

• Temperature equation for conversion.

Kelvin Temperature Scale

• Absolute zero: −273 °C = 0 K.

  • Units: kelvins (K), no degree symbol.

  • No negative temperatures.

  • Same size units as Celsius: 1 K = 1 °C

Calculations and Equations

• TF = 1.8TC + 32

• TC = ewline \frac{TF - 32}{1.8}

Energy

• Energy:

  • Makes objects move or stop moving.

  • Ability to do work.

Kinetic Energy

• Energy of motion (e.g., swimming, water flowing).

Potential Energy

• Stored energy for later use (e.g., water at the top of a dam, compressed spring, chemical bonds).

Heat and Units

• Heat is energy associated with particle movement.

• Measured in joules (J) or calories (cal).

  • Faster particles = greater heat.

• Units for Energy:

  • Joule (J): SI unit of energy.

  • Kilojoule (kJ): 1000 joules.

  • Calorie (cal):

    • Amount of energy to raise the temperature of 1 g of water by 1 °C.

  • Kilocalorie (kcal): 1000 calories.

Energy and Nutrition

• Energy values to calculate kilocalories (kcal) or kilojoules (kJ) for food.

Calorimeters

• Used to measure heat transfer.

• Steel container filled with oxygen and measured water.

• Heat released from burning food sample increases water temperature, which calculates energy value of food.

Energy Values

• Shown as nutritional Calorie (C).

  • 1 Cal = 1000 calories = 1 kcal

• Energy value for 1 g of food in kJ or kcal.

Energy Calculation

• Example:

  • Carbohydrate: 13 g * 4 kcal/g = 52 kcal

  • Fat: 9.0 g * 9 kcal/g = 81 kcal

  • Protein: 9.0 g * 4 kcal/g = 36 kcal

  • Total energy = 52 kcal + 81 kcal + 36 kcal = 169 kcal

Specific Heat

• Specific Heat (SH or c):

  • Different for different substances.

  • Amount of heat to raise the temperature of 1 g of a substance by 1 °C.

  • Units: J/g °C or cal/g °C.

Heat Equation

• q = mc\DeltaT

  • q = heat gained or lost.

  • c = specific heat.

  • \DeltaT = temperature change, \DeltaT = Tf - To (final - initial).

    • Negative q = heat lost; positive q = heat gained.

Calculations

• Can calculate heat lost or gained by measuring mass and temperature change.

Changes of State

• Matter converts from one state to another.

Melting and Freezing

• Melting: solid to liquid at melting point (mp).

• Freezing: liquid to solid at freezing point (fp).

• Water mp/fp: 0 °C.

• Reversible processes.

Evaporation, Boiling, and Condensation

• Evaporation: molecules on surface gain energy and form a gas.

• Condensation: gas molecules lose energy and form a liquid.

• Boiling: all water molecules acquire enough energy to vaporize; bubbles appear throughout the liquid.

Specific Heat Example

A 2.04 g piece of metal at 95.0oC was placed in 35.0 g of H2O at 21.2oC. The temperature of the water increased to 22.1oC. What is the specific heat of the metal?

Assume no heat is exchanged with the surroundings:

-q{hot} = q{cold}

Assume thermal equilibrium is reached: everything present has the same final temperature.

-q{metal} = q{water}

This is a temperature change problem:

q = mc\DeltaT

The specific heat of the metal can be calculated:

c =
ewline \frac{q}{m\DeltaT}

Both m and \deltaT are given directly in the problem, so we only need to find q_{metal} to solve the problem.

Finding q{metal} by first finding q{water}: (-q{hot} = q{cold})

q_{water} = mc\DeltaT

q_{water} = (35
ewline g)(1 \frac{cal}{goC})(22.1oC – 21.2oC)

q_{water} = 31.5
ewline cal

For metal:

c =
ewline \frac{q}{m\DeltaT}

q_{metal} = -31.5
ewline cal

c_{metal} = \frac{(-31.5
ewline cal)}{(2.04
ewline g)(22.1oC – 95.0oC)}

c_{metal} = 0.212 \frac{cal}{goC}